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Relative atomic mass - A r. Just another way of saying how heavy different atoms are compared with the mass of one atom of carbon – 12 (regular carbon!). Mass number. Atomic number. Q. What does the atomic mass number represent?. Number of protons present.

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relative atomic mass a r
Relative atomic mass - Ar

Just another way of saying how heavy different atoms are compared with the mass of one atom of carbon – 12 (regular carbon!)

Mass number

Atomic number

Q. What does the atomic mass number represent?

Number of protons present

A. Total number of protons and neutrons found in an atoms nucleus

another handy trick you can do with the periodic table
Another handy trick you can do with the periodic table

You can work out how many neutrons an element has by subtracting the proton number from the mass number!!

So how many neutrons:

Mass number

Atomic number

relative atomic mass a r3
Relative atomic mass - Ar

Relative atomic mass is easy!! It’s the same value as the mass number – it just sounds scarier!

So what does this tell us about Mg and H?

Table.1. Different elements and their different Ar

relative formula mass m r
Relative formula mass - Mr
  • To find the relative formula mass (Mr) of a compound, you just add together the Ar values for all the atoms in its formula.

Example 1: Find the Mr of carbon monoxide (CO).

The Ar of carbon is 12 and the Ar of oxygen is 16.

So the Mr of carbon monoxide is 12 + 16 = 28.

O

16

8

relative formula mass m r5
Relative formula mass - Mr
  • Example 2:Find the Mr of sodium oxide-Na2O The Ar of sodium is 23 and the Ar of oxygen is 16.

So the Mr of sodium oxide is (23 x 2) + 16 = 62.

slide6

Find the Mr of these:

Carbon dioxide

Sulphur dioxide SO2

Calcium carbonate CaCO3

Sodium hydroxide NaOH

Sulphuric acid H2SO4

Hydrochloric acid HCl

Copper sulphate CuSO4

Magnesium chloride MgCl2

Sodium carbonate Na2CO3

slide7

ATOMIC MASS AND AVERAGE ATOMIC MASS

  • Atomic Mass = used to numerically indicate the mass of an atom in its ground state, it is expressed in the non SI unit of u
  • u = refers to unified atomic mass unit (formerly known as atomic mass unit or amu)
  • 1 amu = 1/12 the mass of carbon-12 atom,
  • therefore the mass of C-12 atom is made equal to 12 amu
slide8

The mass # of an element (periodic table) is the weighted

avg. of allisotopes that exist in nature.

63.55 g/mole

Average Mass of Isotopes

  • Isotopes are naturally occurring.

- abundance of isotope is just as important as mass!

  • Ex...

Natural copper (Cu) consists of 2 isotopes ...

Copper - 63 (mass = 62 .930 g/mole)

69%

Copper - 65 (mass = 64 .930 g/mole)

31%

  • To calculate avg. mass...

mass x abundance for each isotope

Step 1 :

add the two values from step 1 together

Step 2 :

43.42

62 .93 x .69 =

43.42

+

20.13

64 .93 x .31 =

20.13

slide9

The average mass of an element is closest to the isotope

that is mostplentiful in nature.

  • The avg. mass (from P.T.) is closest to 16, therefore,

Oxygen-16 is the isotope that is most abundant in nature.

  • Ex...

Three isotopes of Oxygen:

Oxygen - 16

99 . 759%

Oxygen - 17

0.037%

Oxygen - 18

0.204%

working with weighted averages calculating average atomic mass
Working with Weighted Averages&Calculating Average Atomic Mass

Read page 17 to 18 of your text and make short notes.